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Name _____Mr. Perfect__________________________ Date ____Sp 09__________ 1. Draw Lewis dot structures for the following molecules and include any formal charges on the atoms. (20 pts) a) H3O+ H b) O3 + H O O H c) BH4 - - + O O d) CO H H B H H - C + O 2. Draw the possible resonance structures for N2O4 and include any formal charges on the atoms. (10 pts) O - O + N O - + N O O + N O O O + + N N O O - O - - O O + + N N O O - + N O - 3. Using VSEPR theory, draw the following molecules in their correct molecular geometry and name the molecular geometry. (10 pts) a) SCl4 b) SnCl5 Cl Cl Cl Cl S Cl See-Saw Cl - Sn Cl Cl - Cl Trigonal bipyramidal Chemistry 101 Exam 3 Name _____Mr. Perfect__________________________ Date ____Sp 09__________ 4. Give the expected bond angles for the following molecules and predict what hybrid orbitals are expected for each: (10 pts) a) Bond angles 90 b) Bond angles 90 F F F F S F F Xe F F F sp3d2 F sp3d2 5. Write both the electron configurations and the orbital diagrams for the following elements in their ground states. Also, state if the element is paramagnetic or diamagnetic. (10 pts) a) 18Ar [Ar] [Ar] diamagnetic b) 12Mg [Ne] 3s2 [Ne] [↑↓] 3s diamagnetic c) 28Ni [Ar] 4s23d8 [Ar] [↑↓] [↑↓][↑↓][↑↓][↑ ][↑ ] 4s 3d paramagnetic d) 13Al [Ne] 3s23p1 [Ne] [↑↓] [↑ ][ ][ ] 3s 3p paramagnetic Chemistry 101 Exam 3 Name _____Mr. Perfect__________________________ Date ____Sp 09__________ 6. If a molecule is considered “unstable” with a bond order equal to zero or less, then it is assumed that the molecule does not exist. Using molecular orbital theory, determine if the following molecules are expected to exist. Draw the molecular orbital diagrams for each. (20 pts) {The filling order for the first 2 periods are as follows: 1s *1s 2s *2s 2p 2p *2p *2p} a) H2+ b) He2 ____ 1s* _↑↓_ 1s _↑_ ___ 1s _↑__ H atom ↑↓_ 1s 1s _↑↓ _↑↓_ 1s H atom H2 molecule B.O. = ½ (1-0) = ½ exists 1s* He atom 1 He atom He2 molecule B.O. = ½ (2-2) = 0 does not exist c) B2 d) Be2 ___ 2p ___ ___ _↑_ ___ ___ 2p * 2p* _↑_ ___ ___ 2p ___ _↑↓_ 2p 2s 2s * 2s _↑↓_ _↑↓_ 2s Be Be atom _↑↓_ 2s atom Be2 molecule 2p _↑_ _↑_ _↑↓_ * 2s _↑↓_ _↑↓_ 2s B B atom _↑↓_ 2s atom B2 molecule B.O. = ½ (4-2) = 1 exists Chemistry 101 Exam 3 B.O. = ½ (2-2) = 0 does not exist Name _____Mr. Perfect__________________________ Date ____Sp 09__________ 7. Use partial orbital diagrams to show how the valence shell atomic orbitals of the central atom of the following molecules combine to form a hybrid orbital “Valence Bond Theory”. (Hint: Use VSEPR to predict the hybridization of the central atom.) (10 pts) a) H2O b) BF3 O [↑↓] [↑↓][↑ ][↑ ] atom 2s 2p B [↑↓] [↑ ][ ][ ] atom 2s 2p empty p-orbital [↑ ][↑ ][↑ ] [ ] sp2 p B will now accept an e for each H. [↑↓][↑↓][↑ ][↑ ] sp3 O will now accept an e for each H. 8. Draw C2H4 showing the hybrid orbitals of carbon involved in bonding. Show all of the bonds and bonds. What is the hybridization of carbon in this molecule? (10 pts) bond bond H H H H Each Carbon is Sp2 hybridized bond 9. Extra Credit (5 pts) Consider the following bond lengths: C-O 143 pm C=O 123 pm C=O 109 pm Explain why in the CO32- ion, all three Carbon-Oxygen bonds have identical bond lengths of 136 pm? The bond length is an average of two single bonds plus a double bond due to the following resonance structures. O O C O - O - - O C O O - O Chemistry 101 Exam 3 - - C O